A comparative study on Co(II) removal capacity from water samples by sorption using limestone and nanolimestone

Cobalt(II) ion is a toxic heavy metal ion in industrial wastewater. A trace amount (μgL⁻¹) of cobalt is required and necessary for some organisms as a cofactor for enzymatic activities. However, for most of organisms, concentrations at ppm (5 mg L⁻¹) level are known to be toxic because of the irreversible inhibition of some enzymes by heavy metal ions (Leyssens et al. 2017). In humans, cobalt deficiency has never been reported, but the toxic effects of excess cobalt have been fully described as excess intake will cause acute poisoning by mouth in humans producing gastrointestinal upsets, and chronic absorption may cause abnormalities in skin, heart, blood, or lungs (Payne 1977).

Due to the mobility and toxicity in natural water ecosystems, the presence of Co(II) ions in surface water and groundwater poses a major inorganic contamination problem. To date, there are various technologies for removing heavy metal ions from solution, including filtration, surface complexation, chemical precipitation, ion exchange, adsorption, electrode position, and membrane processing (Al-Qodah 2006).

Most of these processes are unacceptable owing to their high cost, low efficiency, disposal of sludge, and inapplicability to a wide range of pollutants (Quintelas et al. 2009). Adsorption, on the other hand, is one of the most recommended physico-chemical treatment processes that is commonly used and applied for heavy metals' removal from water samples and aqueous solutions. In addition, the adsorption process is well recognized as one of the most efficient methods for removal of heavy metals from their matrices. Adsorption is mainly based on the utilization of solid adsorbents from either organic, inorganic, biological, or low cost materials (Camel 2003). Heavy metal removal via adsorption by organic adsorbents is usually accomplished by the application of polymeric ion-exchangers in which the binding and interaction of metal species with these adsorbents is favored via ion-exchange mechanism, or by application of chelating polymers where the target metal ions are directly attached to the adsorbents via chelating or complex formation mechanism (Gode & Pehlivan 2003). Naturally occurring materials, either modified or unmodified, organic adsorbents such as chitosan and cross-linked carboxymethyl-chitosan (Sun & Wang 2006), polysaccharide-based materials, and lignocellulosic fibers with their surface characteristic functional groups in the form of hydroxyl or carboxyl have also been of great research interest. Biological adsorbents, referred to as biosorbents, are commonly derived from biological components such as bacteria, fungi, and algae, that are characterized and capable of complex formation and/or ion-exchange reactions with metal ions via their functional groups in a process known as biosorption (Abdel-Raouf & Abdul-Raheim 2017).

Removal and extraction of heavy metals based on applications of biosorption approach are commonly performed owing to the major advantages of various biosorbents, such as economical nature, eco-friendly behavior, regeneration for multiple uses, and high selectivity towards different metals (Quintelas et al. 2009; Singh et al. 2017). Low cost adsorbent materials originating from industrial products or wastes are also known as biosorbents and are widely used for the removal of heavy metals from water samples, and these include components of plants, wood, grasses, compost, peat moss, and carbon materials (Romero-Gonzalez et al. 2006).

Inorganic solid adsorbents such as novel silica gel matrix (Elkady et al. 2016), alumina with immobilized 1-nitroso-2-naphthol (Reshetnyak et al. 2012), nano-TiO₂ (Mohammadi & Aliakbarzadeh Karimi 2017), and bentonite adsorbent (Ndifor-Angwafor et al. 2017) are well characterized by their high mechanical properties and strong resistivity to thermal degradation as compared to other biosorbents or organic adsorbents.

In recent years, there has been interest in discovering new products that have minimal environmental impact for restoration or remediation of natural resources (Quintelas et al. 2009).

Limestone (LS) which is produced in large quantities in many countries (among them Egypt) is a low-cost reactive medium that can be used for the retaining of heavy metals and the subsequent clean-up of industrial effluents, leachates, and contaminated ground water (Aziz et al. 2008). Hence, the objective of the present work was to study the possibility of utilizing LS (which is naturally occurring or readily available and cheap) as a sorbent for removing cobalt ions from aqueous solutions and natural waters. The different parameters influencing the adsorption of cobalt ions onto powdered limestone were optimized and the results are presented in this work.

The difference between nanolimestone (NLS) and LS is interpreted as a term of results of XRD and SEM analysis (JSM 6510 LV). Powered XRD (MiniFlex, HyPix-400 MF) studies help to understand the changes occurring in the surface morphologies for the two sorbents. The XRD patterns of LS and NLS (Figure 1) show almost no change in the Lin (counts) for the peak at 2θ which is indicated by the fact that neither has any change of calcium content. In the SEM images of LS and NLS (Figure 2), responding to the morphological changes in the surface of the adsorbents in the coverage of pores with a high porosity for NLS, and the surface deformation, may indicate the incorporation of the metal ion particles into NLS pores.

The limestone (LS), CaCO₃ samples used in this study were obtained from the Al-Mokattam area in Cairo (Egypt) where some private and governorate quarries are located. The samples were crushed and pulverized in the laboratory and those with a mean size of ca. 12.5 mm were used in the experiments.

The samples contained 92% calcite and 3% dolomite (as found by elemental chemical analyzer 2400 Series II, Perkin-Elmer) as has been previously described (Aziz et al. 2008) The remainder of the samples was composed of common minor constituents such as silica, clay, feldspar, pyrite, and sedrite (Bates & Jackson 1980). The samples were dried for 2 h in an oven at 125 °C, packed into stoppered bottles, and stored in a desiccator for future use. Functional groups of LS were characterized through infrared analysis.

The LS spectrum coincided with pure CaCO₃. The surface area and porosity of NLS/LS was measured using the Brunauer, Emmett and Teller (BET) method (Quantachrome TouchWin). LS presented no BET porosity and its measured surface area was 0.50 m² g⁻¹ for limestone and 6.2 m² g⁻¹ for nanolimestone. The pH values of points of zero charge (pHPZC) using a Zetasizer Nano ZSP (Malvern Panalytical) were 9.1 (not aged), 6.2 (aged 60 min) and 8.3 (aged several days), and this agreed with previously reported data (Somasundaran & Goddard 1979). Stirring the LS and NLS sorbent with distilled water (pH = 6.8) for 1 h decreases the suspension pH to 5.8, confirming the positive charge of the LS and NLS surfaces. Also, the concentration of calcium ion in the solution was measured before and after adsorption in order to confirm that cation exchange was involved.

All the solutions were prepared from certified reagent grade chemicals. A cobalt chloride stock solution of 1 molar concentration was prepared and the working solutions were made by diluting the former with doubly distilled water. Aqueous solutions of HNO₃ and NaOH were used for pH adjustments.

Nanolimestone was prepared by dissolving the powdered limestone in about 150 g of concentrated HCl. The obtained calcium chloride solution was mixed with 1.5 g of chitosan, which was dissolved in 3% acetic acid. This mixture was blended with 70 g Na₂CO₃ and slightly heated to complete the reaction. The mixture was kept overnight, then filtered and washed several times with water. Finally, it was calcined in a muffle furnace at 650 °C for 2 hr (Hariharan et al. 2014).

Unless stated otherwise, all batch sorption experiments were conducted at room temperature (ca. 25 °C). Known volumes of Co(II) ions solutions with concentrations ranging from 2 to 100 mg/L were pipetted into quick-fit glass bottles containing 0.1 g of NLS sorbent or 1 g of LS in 100 mL aqueous solution. Since the pH of any of the resulting solutions was about 7.0, no further control was necessary since it was suitable for most adsorption experiments. The resulting solutions were then stirred with a magnetic stirrer at 250 rpm and the samples were taken at fixed time periods (0.0, 0.3, 0.8, 1.0 5.0, 10, 20, 30, 60, and 90 min) in order to study the kinetics of the adsorption process. Preliminary experiments showed that this time length was sufficient for adsorption of Co(II) ions onto NLS or LS. The samples were subsequently filtered off and the residual Co(II) ions concentrations in the filtrate were analyzed using atomic absorption spectroscopy (AAS). Also, the candidate can be reused again in the industrial field (Aliabdo et al. 2014).

The pH value of the solution plays an important role in the sorption of Co(II) ions on LS or NLS. The sorption of Co(II) ions to LS increases from ∼20% to ∼90% in the pH ranging from 4 to 7 and then maintains a high level with increasing pH values (see Figure 3). Similar results were also reported for Cu(II) sorption to TiO₂ (Kim et al. 2003).

The pH values of the solution after sorption equilibrium changes a little to the acidic region, which indicates that H⁺ is released during the sorption process. In aqueous systems, the surface groups of sorbent can be protonated in different extents. Therefore, the concentrations of surface species of the sorbent change under different pH values. With increasing pH, the negative charged groups or deprotonated groups increase and the hydrolysis of Co(II) also increases. As can be seen, below pH 2, the removal of Co(II) ions reaches zero which may be attributed to the complete solubility of LS (consists mainly of CaCO₃ and MgCO₃), thereby hindering the sorption of cobalt ions. Above pH 2, the removal efficiency increases, reaching a maximum value (ca. ∼90%) over the pH range 6–7 followed by a decrease.

The decrease in the removal efficiency at high pH values may be attributed to the fact that the negative species of cobalt, Co(OH)₃⁻ and Co(OH)₄²⁻, are not capable of a combination with the negative surface of LS, as determined by ZPC (pH_ZPC = 6.2 after which the surface is negative). Moreover, this finding was confirmed by stirring NLS or LS with distilled water; the pH of the suspension was always decreased from 6.8 to 5.8. Therefore, pH 7 was recommended throughout all the other experiments.

Adsorption kinetics is one of the most important characteristics representing adsorption efficiency. The adsorption rate of cobalt adsorption on the LS surface, as a function of the initial cobalt concentration, is shown in Figure 4. Due to faster adsorption kinetics with smaller particles (Badruzzaman et al. 2004), the cobalt adsorption was initially rapid, after which it decreased.

The time required to reach equilibrium was 60 min for all cobalt concentrations. The initial rapid adsorption was presumably due to electrostatic attraction. The slow adsorption in the later stage represented a gradual uptake of fluoride at the inner surface by complexation or ion-exchange or due to the decrease in the number of adsorption sites having affinity toward cobalt(II) ions (Sapuan et al. 2013).

In Figure 5, the Morris-Weber curve shows that an initial linear portion may be due to the boundary layer effect (Mohan & Karthikeyan 1997) and a second portion may be due to the intraparticle diffusion effect (Aljeboree et al. 2017). The value of the rate constant for the intrapore diffusion K_d was evaluated as 0.5 (g.g min⁻¹) for LS and 5.6 (g.g min⁻¹) for NLS, which gives an indication about the mobility of the Co(II) ions towards the NLS or LS surface.

The linear plot of Log (q_e−q) vs. t (Figure 6) shows the appropriateness of the above equation and, consequently, the first-order nature of the process involved. The value of K_ads was calculated to be 0.027 min⁻¹ for LS and 0.04 min⁻¹ for NLS.

These results (Figure 7) show that the diffusion of Co(II) ions onto LS pores played a role in the adsorption process (Ali et al. 2015). The value of α constants deduced were 0.11 for LS and 0.21 for NLS, respectively, favored to be (less than 1).

The magnitude of E is useful for estimating the type of sorption reaction. If E is in the range of 8–16 kJ mol⁻¹, the sorption is governed by chemical ion-exchange. In the case of E < 8 kJ mol⁻¹, physical forces may affect the sorption. On the other hand, sorption may be dominated by particle diffusion if E > 16 kJ mol⁻¹ (Sari et al. 2007).

From the results of D-R model simulation (Table 1), E value is 7.6 kJ mol⁻¹ for LS and 7.8 kJ mol⁻¹ for NLS, indicating that sorption is governed by chemical ion-exchange according to the theory of the D-R model or sorption may be considered as physical-chemical adsorption (Özcan et al. 2006).

In order to investigate the effect of temperature on the adsorption of Co(II) ions onto LS, the distribution coefficient, K_d (L.g⁻¹), was calculated at the temperatures 288, 298, 313, and 323 K (Table 2) by using Equation (13).

The K_d value increases with increasing temperature revealing the adsorption of metals onto LS to be endothermic.

Thermodynamic parameters, the enthalpy change (ΔH^o) and the entropy change (ΔS^o), were calculated from the slope and intercept of Equation (13).

The other thermodynamic parameter, Gibbs free energy (ΔG^o) was calculated by Equation (14). For the enthalpy change, the positive ΔH^o showed that the adsorption of Co(II) onto both sorbents is endothermic. For the entropy change (ΔS^o), the positive sign means that the adsorption of Co(II) ions onto both sorbents is a random reaction. Also, the negative sign of the ΔG^o indicates that the adsorption of Co(II) onto both sorbents is feasible and spontaneous thermodyamically. In addition, the reaction proceeded physically, and these results were in good agreement with that obtained from the D-R isotherm.

Table 3 shows the performance for the adsorption of Co(II) onto different adsorbents.

To investigate the applicability of the recommended procedure, a series of experiments was performed to recover 10.0 and 15 mg L⁻¹ of Co(II) ions added to aqueous and some natural water samples using LS and NLS. The adsorption experiments were carried out using clear, filtered, and uncontaminated sample solutions after adjusting their pH values to 7. The results obtained are listed in Table 4; the lower recovery values of Co(II) ions from some water samples could be enhanced by increasing the NLS or LS dose.

The experimental results indicate that NLS and LS can be successfully used for the adsorption of Co(II) ions from aqueous solutions. The equilibrium data well followed the Langmuir and Freundlich models. The value of mean sorption energy, E, obtained from the D-R isotherm indicated that the adsorption of the metals on the LS was feasible and spontaneous. The negative ΔH^o value depicted that the adsorption of Co(II) onto LS was an endothermic process, and the increase in K_d values with increasing temperature also supported this conception. The positive ΔS^o values revealed the randomness of the adsorbed system. On the basis of all results, it can be calculated that limestone (LS) can effectively be used for the removal of cobalt metal cations from different water systems using adsorption method. The nanolimestone presents the major advantage of providing low cost recovery processes making it useful for use in water purification.

The authors extend their appreciation to the Deanship of Scientific Research at King Khalid University for funding this work through Group Research Project under grant number (R.G.P.1/35/38).